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Short description: Inorganic compound with formula H2O3
Structural formula of trioxidane
Trioxidane structure.png
Preferred IUPAC name
Trioxidane (only preselected name)[1]
Other names
Dihydrogen trioxide
Hydrogen trioxide
Dihydroxy ether
3D model (JSmol)
Molar mass 50.013 g·mol−1
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Trioxidane (also systematically named μ-trioxidanediidodihydrogen), also called hydrogen trioxide or dihydrogen trioxide, is an inorganic compound with the chemical formula H[O]3H (also written as [H(μ-O3)H] or [H2O3]). It is one of the unstable hydrogen polyoxides.[2] In aqueous solutions, trioxidane decomposes to form water and singlet oxygen:

Reaction of trioxidane (blue) with water (red) results in decomposition to oxygen and an additional water molecule.

The reverse reaction, the addition of singlet oxygen to water, typically does not occur in part due to the scarcity of singlet oxygen. In biological systems, however, ozone is known to be generated from singlet oxygen, and the presumed mechanism is an antibody-catalyzed production of trioxidane from singlet oxygen.[3]


Trioxidane can be obtained in small, but detectable, amounts in reactions of ozone and hydrogen peroxide, or by the electrolysis of water. Larger quantities have been prepared by the reaction of ozone with organic reducing agents at low temperatures in a variety of organic solvents such as the anthraquinone process, and it is also formed during the decomposition of organic hydrotrioxides (ROOOH).[4] Alternatively, trioxidane can be prepared by reduction of ozone with 1,2-diphenylhydrazine at low temperature. Using a resin-bound version of the latter, relatively pure trioxidane can be isolated as a solution in organic solvent. Preparation of high purity solutions is possible using the methytrioxorhenium(VII) catalyst.[5] In acetone-d6 at −20 °C, the characteristic 1H NMR signal of trioxidane could be observed at a chemical shift of 13.1 ppm.[4] Solutions of hydrogen trioxide in diethyl ether can be safely stored at −20 °C for as long as a week.[5]

The reaction of ozone with hydrogen peroxide is known as the "peroxone process". This mixture has been used for some time for treating groundwater contaminated with organic compounds. The reaction produces H2O3 and H2O5.[6]


In 1970-75, Giguère et al. observed infrared and Raman spectra of dilute aqueous solutions of trioxidane.[2] In 2005, trioxidane was observed experimentally by microwave spectroscopy in a supersonic jet. The molecule exists in a skewed structure, with an oxygen–oxygen–oxygen–hydrogen dihedral angle of 81.8°. The oxygen–oxygen bond lengths of 142.8 picometer are slightly shorter than the 146.4 pm oxygen–oxygen bonds in hydrogen peroxide.[7] Various dimeric and trimeric forms also seem to exist.

There is a trend of increasing gas-phase acidity and corresponding pKa as the number of oxygen atoms in the chain increases in HOnH structures (n=1,2,3).[8]


Trioxidane readily decomposes into water and singlet oxygen, with a half-life of about 16 minutes in organic solvents at room temperature, but only milliseconds in water. It reacts with organic sulfides to form sulfoxides, but little else is known of its reactivity.

Recent research found that trioxidane is the active ingredient responsible for the antimicrobial properties of the well known ozone/hydrogen peroxide mix. Because these two compounds are present in biological systems as well it is argued that an antibody in the human body can generate trioxidane as a powerful oxidant against invading bacteria.[3][9] The source of the compound in biological systems is the reaction between singlet oxygen and water (which proceeds in either direction, of course, according to concentrations), with the singlet oxygen being produced by immune cells.[4][10]

Computational chemistry predicts that more oxygen chain molecules or hydrogen polyoxides exist and that even indefinitely long oxygen chains can exist in a low-temperature gas. With this spectroscopic evidence a search for these type of molecules can start in interstellar space.[7]

See also


  1. Nomenclature of Organic Chemistry : IUPAC Recommendations and Preferred Names 2013 (Blue Book). Cambridge: The Royal Society of Chemistry. 2014. p. 1024. doi:10.1039/9781849733069-FP001. ISBN 978-0-85404-182-4. 
  2. 2.0 2.1 Cerkovnik, J.; Plesničar, B. (2013). "Recent Advances in the Chemistry of Hydrogen Trioxide (HOOOH)". Chem. Rev. 113 (10): 7930–7951. doi:10.1021/cr300512s. PMID 23808683. 
  3. 3.0 3.1 Nyffeler, P.T.; Boyle, N.A.; Eltepu, L.; Wong, C.-H.; Eschenmoser, A.; Lerner, R.A.; Wentworth Jr., P. (2004). "Dihydrogen Trioxide (HOOOH) Is Generated during the Thermal Reaction between Hydrogen Peroxide and Ozone". Angew. Chem. Int. Ed. 43 (35): 4656–4659. doi:10.1002/anie.200460457. PMID 15317003. 
  4. 4.0 4.1 4.2 Plesničar, B. (2005). "Progress in the Chemistry of Dihydrogen Trioxide (HOOOH)". Acta Chim. Slov. 52: 1–12. 
  5. 5.0 5.1 Strle, G.; Cerkovnik, J. (2015), "A Simple and Efficient Preparation of High‐Purity Hydrogen Trioxide (HOOOH)", Angew. Chem. Int. Ed. 54 (34): 9917–9920, doi:10.1002/anie.201504084, PMID 26234421 
  6. Xu, X.; Goddard, W.A. (2002). "Nonlinear partial differential equations and applications: Peroxone chemistry: Formation of H2O3 and ring-(HO2)(HO3) from O3/H2O2". PNAS 99 (24): 15308–15312. doi:10.1073/pnas.202596799. PMID 12438699. 
  7. 7.0 7.1 Suma, K.; Sumiyoshi, Y.; Endo, Y. (2005). "The Rotational Spectrum and Structure of HOOOH". J. Am. Chem. Soc. 127 (43): 14998–14999. doi:10.1021/ja0556530. PMID 16248618. 
  8. Plesničar, Božo (2005). "Progress in the Chemistry of Dihydrogen Trioxide (HOOOH)". Acta Chim. Slov. 52: 1–12. 
  9. A Time-Honored Chemical Reaction Generates an Unexpected Product, News & Views, September 13, 2004
  10. Hoffmann, R. (2004). "The Story of O". Am. Sci. 92: 23. doi:10.1511/2004.1.23.